Li_1.6AlCl_3.4S_0.6: a low-cost and high-performance solid electrolyte for solid-state batteries

Abstract

Solid electrolytes (SEs) are crucial for advancing next-generation rechargeable battery technologies, but their commercial viability is partially limited by expensive precursors, unscalable synthesis, or low ionic conductivity. Lithium tetrahaloaluminates offer an economical option but exhibit low Li⁺ conductivities with high activation energy barriers. This study reports the synthesis of lithium aluminum chalcohalide (Li_1.6AlCl_3.4S_0.6) using inexpensive precursors via one-step mechanochemical milling. The resulting Cl–S mixed-anion sublattice significantly improves the ionic conductivity from 0.008 mS cm⁻¹ for LiAlCl₄ to 0.18 mS cm⁻¹ for Li_1.6AlCl_3.4S_0.6 at 25 °C. Structural refinement of the high-resolution XRD patterns and ⁶Li magic-angle-spinning (MAS) NMR quantitative analysis reveals the formation of tetrahedrally-coordinated, face- and edge-shared LiCl_xS_y octahedra that facilitate 3D Li⁺ transport. Ab initio molecular dynamics (AIMD) simulations on Li_1.6AlCl_3.4S_0.6 support an enhanced 3D network for Li⁺ migration with increased diffusivity. All-solid-state battery (ASSB) half-cells using Li_1.6AlCl_3.4S_0.6 exhibit high-rate and long-term stable cycling performance. This work highlights the potential of Li_1.6AlCl_3.4S_0.6 as a cost-effective and high-performance SE for ASSBs.

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Introduction

All-solid-state batteries (ASSBs) are rapidly gaining traction due to their high energy density and enhanced safety features, positioning them as strong contenders to replace traditional liquid-electrolyte-based lithium-ion batteries (LIBs) in portable electronic devices and electric vehicles.^1,2 The practical commercialization of ASSBs hinges on discovering SEs that satisfy the cost and performance metrics. In recent years, ternary halides such as LiAlCl₄ and NaAlCl₄ have received significant attention owing to their low cost and fast ion transport when in solution or as a melt.^3–6

Aluminum is the third most abundant element in the earth's crust,⁷ which makes it a desirable lost-cost component in SEs for ASSBs. However, despite recent advancements in enhancing Li⁺ dynamics within the lithium tetrahaloaluminate class of solid electrolytes, their ionic conductivity remains notably low, thereby restricting their practical applications in ASSBs. Recent computational studies have demonstrated that stabilizing a metastable phase of LiAlCl₄ can force Li-ions to occupy interstitial sites, thereby reducing the activation energy barrier for ion transport. Specifically, mechanochemical synthesis has induced structural disorder in LiAlCl₄, resulting in a small fraction (2.5%) of Li-ions occupying interstitial tetrahedral sites.⁸ This has significantly improved the ionic conductivity of LiAlCl₄ to 0.02 mS cm⁻¹ at 25 °C, compared to the previous value of 0.001 mS cm⁻¹.⁹ However, this improvement is still insufficient for commercial applications in ASSBs, which require conductivities greater than 0.1 mS cm⁻¹ at 25 °C.

A synthesis approach that does not require high-temperature heating is highly attractive for fabricating commercially viable SEs. For example, by subjecting LiAlBr₄ to mechanical stress, a moderate ionic conductivity of 0.033 mS cm⁻¹ at room temperature has been reported.³ Li⁺ transport within this material is facilitated via octahedral and tetrahedral interstitial sites, which are intrinsic to lithium tetrahaloaluminates.^3,8 Previous studies have shown that for some Li-SEs, e.g., Li₆PS₅Cl, Li₃PS₄, and LiAlCl₄, ionic conductivity can be enhanced by introducing local disorder within the anion sublattice.^10–17 Moreover, studies have reported connecting non-accessible ion-conducting channels by a random distribution of cations in a host structure via local disorder.¹⁸ To further improve the ionic conductivity, local disorder on the atomic scale can be implemented.¹⁹

Herein, we demonstrate an effective approach for improving Li⁺ conduction in LiAlCl₄ by mixing inexpensive and poor-conducting Li₂S and AlCl₃. An ionic conductivity of 0.18 mS cm⁻¹ is achieved at 25 °C for Li_1.6AlCl_3.4S_0.6 compared to 0.008 mS cm⁻¹ for LiAlCl₄. Powder X-ray diffraction and solid-state nuclear magnetic resonance (NMR) spectroscopy, combined with Ab initio molecular dynamics (AIMD) simulations, were utilized to determine the structural origin of the improved ionic conductivity in Li_1.6AlCl_3.4S_0.6. Galvanostatic measurements of Li_1.6AlCl_3.4S_0.6-containing ASSB half-cells show an improved rate performance and long-term cycling stability compared with the control cell utilizing LiAlCl₄.

Experimental

Material synthesis

LiCl (Sigma-Aldrich) was dried at 200 °C for 12 hours under a dynamic vacuum before being stored in an argon-filled glovebox. Anhydrous Li₂S (Alfa Aesar) and ultra-dry AlCl₃ (Alfa Aesar) were used as received and handled under argon. Stoichiometric amounts of precursors were ground using a mortar/pestle for 5 minutes inside an argon-filled glovebox. The hand-milled powder was transferred into a ZrO₂ jar containing two 10 mm balls as milling media. After the jar was vacuum sealed, mechanochemical mixing was performed using a SPEX 8000M MIXER/MILL (SPEX®SamplePrep, USA) for 20 hours. Afterward, the ball-milled powder, typically ∼130 mg, was pressed into a pellet of 8 mm in diameter under ∼400 MPa inside an argon-filled Mbraun glovebox. The resulting pellet had a thickness of ∼1.2 mm and appeared shiny white.

Powder X-ray diffraction. The as-milled samples were packed in a zero-background sample holder. KAPTON® film (DUPONT™, USA) was used to seal the samples to prevent exposure to humid air. Powder X-ray Diffraction (PXRD) was performed using a RIGAKU® Smartlab powder diffractometer with Bragg-Brentano geometry at a voltage of 45 kV and current of 40 mA with Cu-Ka radiation (a = 1.540 562 Å). The data was collected in the 2θ range of 10–80° at a step size of 0.03° for 30 minutes.

Synchrotron X-ray diffraction. Synchrotron X-ray diffraction (SXRD) measurements were carried out in the transmission mode at the 17-BM-B beamline, APS, at Argonne National Lab (ANL), Illinois. The exact X-ray wavelength was refined to 0.24117 Å. The sample was loaded inside a special glass capillary, and the holder was moved up and down during tests to ensure uniformity of the measured results.

Rietveld refinement. Rietveld refinement of the lab and synchrotron PXRD data was performed using the GSAS-II software. Structural analysis of the synchrotron PXRD data on LiAlCl₄ and Li₂AlCl₃S reveals a monoclinic phase belonging to the P2₁/c space group of LiAlCl₄ (ICSD – 35275). Sulfur occupancy was tested on all the chlorine sites, and the result was only considered accurate when sulfur occupancy was greater than 1% on the tested site. Lithium and aluminum occupancy were tested for the counter-cation sites of each other to find the most probable structure. Atomic parameters for the site were fixed for the substituents (S, Cl, etc.).

Computational approach. All density functional theory (DFT) energy calculations and ab initio molecular dynamics (AIMD) simulations were carried out in the Vienna ab initio simulation package (VASP).²⁰ The projector-augmented-wave (PAW) approach was used.²¹ Perdew–Burke–Ernzerhof generalized-gradient approximation (GGA-PBE) was chosen as the exchange-correlation functional using the latest PAW potential files available in VASP.²² Python Materials Genomics (Pymatgen) package²³ was used to optimize the structures of Li₂AlCl₃S. 10 supercells with different local environments were generated based on the 2 × 2 × 1 supercell of the LACS obtained from the high-resolution XRD structure refinement. Geometry optimization of the generated supercells was carried out using DFT calculations. The AIMD simulations²⁴ were based on the canonical ensemble for over 80 ps with a time step of 2 fs. The temperature was initialized at 100 K and elevated to the target value for the simulations.

Solid-state NMR. ⁶Li and ⁷Li NMR experiments were performed using a Bruker Advance-III 500 spectrometer at Larmor frequencies of 73.6 MHz and 194.4 MHz for ⁶Li and ⁷Li, respectively. The magic-angle-spinning (MAS) rate was 24 kHz. Single-pulse MAS ⁶Li and ⁷Li NMR experiments were performed using π/2 pulse lengths of 3.30 μs and 2.90 μs, respectively. For Li₂AlCl₃S, the recycle delays were 500 s for ⁶Li and 80 s for ⁷Li, while recycle delays of 500 s for ⁶Li and 90 s for ⁷Li was utilized for LiAlCl₄. ⁷Li NMR spectra were calibrated with LiCl(s) at −1.1 ppm, and ⁷Li T₁ relaxation time was measured with an inversion-recovery pulse sequence.

Variable-temperature ⁷Li T₁ relaxation NMR experiments were performed using a Bruker Avance-I 300 MHz Spectrometer from 25 to 70 °C. An inversion recovery pulse sequence with a π/2 pulse length of 2.63 μs was utilized. The ⁷Li Larmor frequency was 116.6 MHz. Sample powders were packed into 4 mm ZrO₂ rotors under Argon and spun at a MAS rate of 10 kHz.

A ⁶Li|SE|⁶Li symmetrical cell was assembled in an argon-filled glovebox for the ⁶Li → ⁷Li tracer exchange experiments. The cell was then subjected to galvanostatic cycling for 3 days at a current density of 5 μA cm⁻² to drive the diffusion of ⁶Li⁺ ions from the ⁶Li foil into the Li₂AlCl₃S pellet. Following the galvanostatic cycling, ⁶Li MAS NMR experiments were performed on the cycled Li₂AlCl₃S pellet using the Bruker Advance-III 500 spectrometer at a spinning rate of 24 kHz, using the same parameters described above.

Electrochemical impedance spectroscopy (EIS). The samples were pressed in a mold of 8 mm diameter to make 1.3 mm thick pellets, which were sandwiched between indium of diameter 6 mm (about 0.24 in) followed by stainless steel plungers as ion-blocking electrodes. EIS measurement was carried out on a Biologic SP-300 electrochemical analyzer within a frequency range from 7 MHz to 1 Hz using a voltage of 10 mV. The conductivities are calculated using resistance obtained by fitting the Nyquist plots using an equivalent circuit model. Variable-temperature EIS characterization was performed in the CSZ microclimate chamber from −20 °C to 70 °C using a Biologic SP-300, and Arrhenius-type plots were used to calculate the activation energies and Arrhenius prefactors.

Direct current (DC) polarization. The DC polarization method was used to measure the electronic conductivity.²⁵ In-house-built split cells (diameter = 8 mm) using PEEK insulating cylinder and stainless-steel plungers as current collectors and indium foils (∼6 mm diameter) as ion-blocking electrodes were used.

Cyclic voltammetry (CV), linear sweep voltammetry (LSV), and galvanostatic cycling of ASSB half-cells. ASSB half-cells were assembled using pressure cells constructed in-house, using a PEEK casing of 10 mm diameter and stainless-steel plungers. For CV measurements, the initial steps involved pressing 100 mg of LiAlCl₄ (or Li₂AlCl₃S) at 300 MPa for 10 s. Then, roughly 10 mg of the manually mixed 3SE : carbon black (C: Super P) composite was evenly spread and pressed at 300 MPa for 10 s. On the opposite side of the LiAlCl₄ (or Li₂AlCl₃S) pellet, a piece of indium (In) foil measuring 5/16 inch in diameter and 0.1 mm in thickness, with an approximate weight of 32 mg was attached. Subsequently, lithium foil with a 3/16-inch diameter and weighing around 1 mg was pressed onto the In foil to form Li–In and used as the counter electrode, giving a final cell setup of Li–In|SE|3SE:C. With the cell sealed using vacuum grease, it was subjected to electrochemical cycling under an estimated stack pressure of approximately 30 MPa at 22 °C. CV and LSV measurements were conducted with a scan rate of 0.2 mV s⁻¹ from 0 to 4 V vs. Li–In and 0 to 4.5 V respectively.

To prepare the composite cathode (or catholyte) for galvanostatic cycling, spherical TiS₂ particles of ∼74 microns were purchased from Sigma-Aldrich. It was dried at 200 °C for 12 hours, then subjected to ball milling for 5 hours at 300 rpm to reduce particle size. Subsequently, LiAlCl₄ or Li₂AlCl₃S was combined with TiS₂ at a TiS₂ : SE mass ratio of 1 : 2 and ground together using a mortar and pestle for 10 minutes. Li₆PS₅Cl, synthesized following the established method by Patel et al.,¹³ was pressed into pellets at 300 MPa for 10 seconds as the separator. For the half-cell assembly, 12 mg of the catholyte was evenly spread onto one side of the Li₆PS₅Cl pellet, achieving an aerial loading of approximately 1.25 mA h cm⁻², followed by further pressing at 300 MPa for 10 seconds. A Li–In alloy foil was affixed to the opposite side of the Li₆PS₅Cl pellet to assemble the Li–In|Li₆PS₅Cl|2SE:TiS₂ (SE: LiAlCl₄ or Li₂AlCl₃S) half cells. Finally, the cells were sealed with vacuum grease and subjected to controlled cycling conditions at 22 °C with a stack pressure of ∼30 MPa, within a voltage window of 1–2.5 V vs. Li–In. For rate performance evaluations, the cells underwent cycling for 5 cycles at each of the following rates: 0.1C, 0.2C, 0.5C, 1C, and 2C, with C representing the charge–discharge rate. Correspondingly, these rates translate to current densities of 0.14 mA cm⁻², 0.28 mA cm⁻², 0.70 mA cm⁻², 1.40 mA cm⁻², and 2.80 mA cm⁻², respectively. Subsequently, long-term stability testing was conducted over 175 cycles at 0.2C.

Results and discussion

X-ray diffraction and structure

X-ray diffraction (XRD) was employed to investigate the long-range structure of the synthesized solid electrolytes. The powder XRD confirms the presence of a monoclinic LiAlCl₄ phase (Fig. 1). With sulfur substitution, the long-range monoclinic structure is maintained for Li₂AlCl₃S. However, the weak and diffuse diffraction pattern of the as-milled Li₂AlCl₃S suggests a decrease in the crystallinity of the material. Rietveld refinement was performed against high-resolution XRD data using GSAS-II. Fig. 1(c)–(f) shows the diffraction patterns and the refined structures. For LiAlCl₄, the crystalline phase was refined with a monoclinic structure type in the P2₁/c space group. The structure is built from distorted LiX₆ octahedra and AlCl₄⁻ tetrahedra. Two LiCl₆⁵⁻ octahedra are edge-shared to form Li₂Cl₁₀⁸⁻ dimers.¹ AlCl₄⁻ tetrahedra are isolated from one another. All the atoms in the structure occupy only the 4e Wyckoff positions, leaving all other sites, i.e., 2a, 2b, 2c, and 2d Wyckoff positions, vacant (see Fig. 1(e)). The arrangement of atoms is characterized by a slightly distorted hexagonal cubic packed (hcp) Cl⁻ sublattice, where Li⁺ and Al³⁺ fill octahedral and tetrahedral interstices, respectively.¹ In addition, each AlCl₄ tetrahedron is linked to one Li–Cl dimer via two edges and to two other dimers through one corner each (Fig. 1(e)).

Fig. 1 Long-range structures of LiAlCl₄ and Li₂AlCl₃S, determined using lab-X-ray diffraction and high-resolution X-ray diffraction. (a) Lab X-ray diffraction patterns of the as-milled LiAlCl₄ and Li₂AlCl₃S. The ICSD patterns of precursors and LiAlCl₄ monoclinic phase (P2₁/c space group) are shown as references. (b) The magnified view of the selected 2θ ranges in (a). (c) and (d) High-resolution X-ray diffraction patterns and the corresponding Rietveld refinement of LiAlCl₄ and Li₂AlCl₃S, respectively. (e) and (g) Monoclinic structures of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6 with the P2₁/c space group, respectively, obtained from refined high-resolution XRD patterns. (f) The structure of Li_1.6AlCl_3.4S_0.6 viewed along the a-axis showing tetrahedral voids face-shared with Li-octahedra. (h) The visualization of partially occupied octahedral lithium at Wyckoff 2b position face-shared with highly distorted octahedral sites at Wyckoff 4e positions. (i) Possible lithium migration pathways.

Similarly, Li₂AlCl₃S crystallizes in the P2₁/c space group. However, different from LiAlCl₄, the unit cell consists of three octahedrally coordinated cation sites – Li1 at 4e, Li2 at 2a, and Li3 at the 2a Wyckoff positions; and tetrahedrally coordinated Al at the 4e Wyckoff position. Sulfur and chlorine atoms co-occupy the 4e anionic site, yielding a disordered anion sublattice. Along the c-direction, the structure exhibits three distinct cation layers (Fig. 1(g)). In the first layer, Li2 octahedra (LiCl₆)⁵⁻ edge-share with Li1 (LiCl_5.4S_0.6)^5.6− and corner-share with the tetrahedrally coordinated Al. In the second layer, Li3 octahedra (LiCl_4.8S_1.2)^6.2− face-share with Li1 (Fig. 1(h)), and both Li1 and Li3 octahedra share corners with the tetrahedrally coordinated Al. In the third layer, the Al sites face-share with Li3 and edge-share with Li1 and Li2 octahedra. In addition, the edge-sharing Li1, Li2, and Li3 octahedra produce tetrahedral voids (Fig. 1(f)). The complete refinement parameters for LiAlCl₄ and Li₂AlCl₃S are provided in Tables S1 to S3.† The refinement reveals the actual composition of the nominal Li₂AlCl₃S is Li_1.6AlCl_3.4S_0.6, which exhibits mixed Cl⁻/S²⁻ site occupancies in the structure. Consistent with the refined composition, a trace amount of Li₂S is identified in the ball-milled sample. The refined composition, Li_1.6AlCl_3.4S_0.6, is adopted hereafter.

Solid-state NMR

To understand the local structures of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6, ⁶Li MAS NMR experiments were performed. As shown in Fig. 2, the ⁶Li MAS NMR spectrum of LiAlCl₄ reveals a major resonance at −0.92 ppm assigned to octahedrally coordinated Li1. This is consistent with findings from a previous study.⁸ In addition, a minor resonance at −1.08 ppm is assigned to LiCl.²⁶ After incorporating S²⁻ in the parent material, the ⁶Li NMR spectra of Li_1.6AlCl_3.4S_0.6 is characterized by five distinct resonances at −0.89, −1.24, −0.07, 1.48, and 2.21 ppm – indicating changes in the local Li environment. The minor peak observed at 2.21 ppm corresponds to the Li₂S impurity.²⁷ The phase quantification is given in Table S4.† The resonances at −0.89, −1.24, and −0.07 ppm are assigned to the octahedrally coordinated Li1, Li2, and Li3 sites, respectively – in agreement with the results from structural refinement. In addition, there is an extra peak at 1.48 ppm, likely from the Li_1.66S_0.66Cl_0.34 phase identified with the Rietveld refinement of the high-resolution diffraction pattern (Fig. S1†).

Fig. 2 ⁶Li MAS NMR spectra of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6.

⁷Li spin–lattice relaxation time (T₁) is a useful indicator of ion dynamics.^11,28 According to the Bloembergen, Purcell, and Pound (BPP) model, T₁ relaxation time is a function of motional rate (τ_c⁻¹),²⁹ where γ is the gyromagnetic ratio, ħ is the reduced Planck's constant, r₀ is the interatomic distance, ω₀ = γB₀ is the Larmor frequency, and B₀ is the external magnetic field strength.

In the fast-motion regime (ω₀τ_c ≪ 1), T₁ increases with increasing motional rate, while in the slow-motion regime (ω₀τ_c ≫ 1), T₁ decreases with increasing motional rate. In addition, a resonance can lie in the intermediate region where ω₀τ_c ≈ 1.²⁹ Variable-temperature ⁷Li NMR T₁ relaxation data of Li_1.6AlCl_3.4S_0.6 in Fig. S2† reveals a decrease in T₁ relaxation time with increasing temperature and, thus, suggesting Li⁺ dynamics in Li_1.6AlCl_3.4S_0.6 lie in the slow-motion regime (ω₀τ_c ≪ 1). Therefore, a shorter T₁ value will correlate with faster ion mobility. As presented in Table 1, the ⁷Li T₁ relaxation time significantly decreases from LiAlCl₄ to Li_1.6AlCl_3.4S_0.6, suggesting enhanced Li⁺ mobility³⁰ with sulfur incorporation.

Table 1 ⁷Li spin–lattice relaxation time (T₁) of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6

Sample	^7Li T1 [s]
LiAlCl4	5.2
Li1.6AlCl3.4S0.6	3.1

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Ion transport pathways determined by tracer-exchange NMR

To directly probe the Li⁺ transport pathways in Li_1.6AlCl_3.4S_0.6, tracer-exchange NMR is employed.³¹ By identifying and quantifying ⁶Li⁺ → ⁷Li⁺ exchange under an applied biased potential, the preferential pathway utilized by Li⁺ ions for migration is directly mapped out – facilitating the identification of active sites for Li⁺ transport.^31–34 The experimental configuration involves sandwiching Li_1.6AlCl_3.4S_0.6 pellet between two ⁶Li foils. An externally applied potential gradient establishes a driving force for ⁶Li⁺ ions in the ⁶Li foils to move toward and exchange with ⁷Li⁺ in the Li_1.6AlCl_3.4S_0.6 pellet. Consequently, the preferential Li⁺ transport pathways are selectively enriched with ⁶Li⁺ ions. The ⁶Li NMR of the pristine and tracer-exchanged Li_1.6AlCl_3.4S_0.6 pellet is displayed in Fig. S3a,† and the quantification of the Li sites before and after the tracer-exchange experiment is presented in Fig. S3b.† The ⁶Li NMR spectra reveal changes in the relative intensities of resonances assigned to distinct Li⁺ environments. Notably, a significant increase in the intensity of the Li1 and Li2 resonances is observed, suggesting the major involvement of Li1 and Li2 in ion conduction. In addition, the Li3 resonance shows a small intensity enhancement after cycling. The enhancement of these resonances suggests that Li1, Li2, and Li3 all participate in Li⁺-ion transport within the Li_1.6AlCl_3.4S_0.6 solid electrolyte.

AIMD simulations

To further understand the effect of Cl–S anion mixing on the Li⁺ density distribution and diffusion, AIMD simulations are employed for LiAlCl₄ and Li_1.6AlCl_3.4S_0.6 in a 2 × 2 × 1 cell. The mean square displacements (MSD) of Li⁺ (Fig. 3(a) and (b)) and distribution probability (Fig. 3(c) and (d)) for Li at 900 K were calculated. The MSD plots demonstrate that Li⁺ can diffuse in all three directions, with higher diffusion observed along the b direction in LiAlCl₄. However, Li⁺ diffusion along the a and c directions increases significantly with Cl–S anion mixing, resulting in comparable MSDs in all three directions in Li_1.6AlCl_3.4S_0.6. This transformation indicates that Li_1.6AlCl_3.4S_0.6 is a 3D ion conductor with a uniform ion transport network in all three dimensions. Furthermore, the overall MSD of Li⁺ shows a 50% increase upon Cl–S anion diversification compared to LiAlCl₄. The AIMD simulation of Li⁺ trajectories (i.e., Li⁺ probability density) in LiAlCl₄ shows a localized “cage-like” Li⁺ diffusion pattern with few interstitial jumps, indicating limited long-range Li⁺ migration. In comparison, Li_1.6AlCl_3.4S_0.6 exhibits a delocalized Li⁺ diffusion network, suggesting improved macroscopic Li⁺ migration and, consequently, enhanced Li⁺ conduction in Li_1.6AlCl_3.4S_0.6.

Fig. 3 Mean square displacement (MSD) of Li⁺ in (a) LiAlCl₄ and (b) Li_1.6AlCl_3.4S_0.6 generated from AIMD simulations. Li⁺ (yellow) probability density map of LiAlCl₄ (c) and Li_1.6AlCl_3.4S_0.6 (d) in a 2 × 2 × 1 cell based on AIMD simulations at 900 K.

Electrochemical properties

To examine ion transport properties of all prepared SEs, variable-temperature electrochemical impedance spectroscopy (EIS) was employed, and the corresponding Nyquist plots at 25 °C are presented in Fig. 4(a) with a corresponding exemplary equivalent circuit fitting shown in Fig. 4(b). The results from the EIS analysis are given in Table 2. Based on fitted resistances from the Nyquist plot using equivalent circuit modeling, the conductivities of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6 are 0.008 mS cm⁻¹ and 0.18 mS cm⁻¹, respectively (Fig. 4(a)). In addition, an ionic conductivity of 0.12 mS cm⁻¹ and 0.11 mS cm⁻¹ were obtained at x = 0.5 and x = 0.8 in Li_1+xAlCl_4−xS_x (Fig. S4†). The lower ionic conductivity observed at x = 0.5 is likely due to insufficient sulfur substitution to significantly perturb the lattice and create additional favorable Li⁺ migration pathways. Conversely, at x = 0.8, the solid-solution limit is reached yielding a high fraction of Li₂S impurity, which likely contributes to the lower ionic conductivity. Therefore, the optimal composition for the highest ionic conductivity is Li_1.6AlCl_3.4S_0.6. To confirm the conductivities are due to ionic transport, the DC polarization measurements²⁵ were performed, and the plots for LiAlCl₄ and Li_1.6AlCl_3.4S_0.6 are shown in Fig. 4(c). The current value upon plateauing corresponds to electronic conductivities of 3.49 × 10⁻⁸ S cm⁻¹ and 5.43 × 10⁻⁸ S cm⁻¹ for LiAlCl₄ and Li_1.6AlCl_3.4S_0.6, respectively, thereby confirming negligible electronic contribution to the measured total conductivities of the samples. The ionic conductivity of Li_1.6AlCl_3.4S_0.6 is approximately twenty-fold greater than LiAlCl₄ at room temperature. The significant increase in ionic conductivity of Li_1.6AlCl_3.4S_0.6 can be attributed to the local disorder and “Li-stuffing” of the Li3-octahedra that connect with neighboring Li1/Li2 octahedra and the substitution of Cl with more polarizable S anions. The energy barrier for Li-ion transport (E_a) was calculated from the variable-temperature EIS measurements between 0 °C to 70 °C, and representative Nyquist plots are shown in Fig. S5.† The Arrhenius-type plot of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6 are shown in Fig. 4(d). A decrease in E_a from 0.51 eV to 0.44 eV is observed from LiAlCl₄ to Li_1.6AlCl_3.4S_0.6. This decrease in activation energy correlates with the increase in ionic conductivity.³⁵ This trend agrees with the energy barrier obtained from the bond valence site energy (BVSE) calculation (Fig. S6†).

Fig. 4 Charge transport properties of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6. (a) Nyquist plots of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6. (b) Exemplary equivalent circuit fitting of the Nyquist plot for Li_1.6AlCl_3.4S_0.6 measured at 25 °C. (c) Electronic conductivities were determined using DC polarization, and (d) Arrhenius plots and activation energies for ion transport in LiAlCl₄ and Li_1.6AlCl_3.4S_0.6.

Table 2 DC ionic conductivity at 25 °C, electronic conductivity at 25 °C, activation energy, and Arrhenius prefactor of LiAlCl₄ and Li_1.6AlCl_3.4S_0.6

Composition	σ DC, 25°C [S cm^−1]	σ e, 25°C [S cm^−1]	E a [eV]	Log(σ0) [S cm^−1 K]
LiAlCl4	8.65 × 10^−6	3.49 × 10^−8	0.51	6.55
Li1.6AlCl3.4S0.6	1.80 × 10^−4	5.43 × 10^−8	0.44	6.07

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Cyclic voltammetry and galvanostatic cycling of ASSB half-cells

Conventionally, CV measurements were done using stainless steel as the blocking electrode, which fails to accurately measure the oxidation and reduction current of SEs due to the limited electrical contact area of the SE and the planar ion-blocking electrode.^11,36 To overcome this and estimate the oxidation–reduction reactions of the SE, we performed CV utilizing a 3SE : C (mass ratio) composite cathode in the Li–In|SE|3SE : C half-cell setup.^37–39 In this setup, carbon serves as an electronic conductive medium in the composite cathode, enabling increased SE surface area in electrical contact with the electrode and consequential detection of degradation current.^{11,17,36–42} Fig. S7† shows the voltammograms of Li_1.6AlCl_3.4S_0.6 with a scanning window of 0–4 V vs. Li–In. The first cathodic peak starts at the voltage of 1.03 V vs. Li–In. This voltage is assigned to the electrolyte reduction to form Li₂S and AlCl₃.^11,17,43 The CV for Li_1.6AlCl_3.4S_0.6 indicates the onset oxidation voltage of 2.4 V vs. Li–In. This is reasonable as most of the reported sulfide solid electrolytes have a lower stability window than halides (LiAlCl₄).^44,45 So, the electrochemical stability window of the prepared Li_1.6AlCl_3.4S_0.6 is in the range of 1.03–2.40 V vs. Li–In, corresponding to the 1.63–3.00 V vs. Li/Li⁺. We have also performed LSV measurement (Fig. S8†), which shows the same electrochemical stability window of 1.63–3.00 V vs. Li/Li⁺

Titanium disulfide (TiS₂) was employed as the cathode active material (CAM) in the half-cell configurations for electrochemical evaluation. Li–In|Li₆PS₅Cl|2(LiAlCl₄):TiS₂ and Li–In|Li₆PS₅Cl|2(Li_1.6AlCl_3.4S_0.6):TiS₂ cells were fabricated according to previous studies.^11,17 A Li–In anode was employed for enhanced stability against SEs and limited dendrite formation through micropores – reducing the risk of short circuits.⁴⁶ Li₆PS₅Cl was utilized as the separator due to its high ionic conductivity and stability against Li metal. The electrochemical performance of these half-cells was assessed through a series of rate capability tests, spanning charging/discharging currents from 0.1C to 2C conducted under galvanostatic conditions at ∼22 °C. Each rate was applied over five cycles (0.1C ∼0.14 mA cm⁻², 0.2C ∼0.28 mA cm⁻², 0.5C ∼0.70 mA cm⁻², 1C ∼1.40 mA cm⁻², and 2C ∼2.80 mA cm⁻²) followed by 125 cycles at 0.2C. A theoretical capacity of 239 mA h g⁻¹ for TiS₂ was used to calculate the charge–discharge rates.

Fig. 5(a) illustrates the variation in specific capacity with cycle number for the half-cell configurations, while Fig. 5(b) and (c) display the voltage profiles for select cycles. Initially, the cell incorporating LiAlCl₄ demonstrated a discharge capacity of approximately 273 mA h g⁻¹ and a charge capacity of 213 mA h g⁻¹, whereas the Li_1.6AlCl_3.4S_0.6-containing cell exhibited a comparable initial discharge capacity of 272 mA h g⁻¹ and a higher charge capacity of 248 mA h g⁻¹ (Fig. 5(a) and (b)). Upon the second discharge, the capacity for the LiAlCl₄-based cell declined to 208 mA h g⁻¹, whereas the Li_1.6AlCl_3.4S_0.6-based cell showed a reduced yet more stable capacity of 240 mA h g⁻¹, eventually stabilizing around 239 mA h g⁻¹ (Fig. 5(a) and (b)). This is likely due to the formation of a stable solid–electrolyte interface (SEI) during the first cycle. The lower capacity observed in the LiAlCl₄-based cell is primarily due to its low ionic conductivity of the solid electrolyte (SE), which restricts the efficient utilization of the active material – leading to a lower capacity.⁴⁷ Interestingly, initial capacities recorded for both cells surpassed the theoretical capacity of TiS₂, (239 mA h g⁻¹). This is attributable to the unidentified reversible redox behavior of the SE in addition to the Ti^3+/4+ redox (Fig. S9†).^48,49

Fig. 5 Galvanostatic cycling of 2SE:TiS₂ catholyte with argyrodite Li₆PS₅Cl as the separator. (a) Rate performance of Li–In|Li₆PS₅Cl|2SE:TiS₂ cell (SE = LiAlCl₄ or Li_1.6AlCl_3.4S_0.6) followed by long-term cycling at C/5, where C = 239 mA h g⁻¹. Voltage profile of cycle 1 along with the 2nd cycles of each C-rate for a (b) Li–In|Li₆PS₅Cl|2(Li_1.6AlCl_3.4S_0.6):TiS₂ cell and (c) Li–In| Li₆PS₅Cl|2(LiAlCl₄):TiS₂ cell.

The Li_1.6AlCl_3.4S_0.6-containing cell exhibited a capacity of 180 mA h g⁻¹ even at a high discharge rate of 2C, indicative of exceptional rate performance in contrast to the significantly reduced capacity of 37 mA h g⁻¹ observed at a similar rate for the LiAlCl₄-based cells. Upon returning to 0.2C after 26 cycles, both cell configurations demonstrated remarkable stability over 200 cycles. Furthermore, these cells maintained a high coulombic efficiency exceeding 99% throughout this extended cycle period (Fig. 5(a)). Between the 27th and 200th cycles, the Li_1.6AlCl_3.4S_0.6-based cell demonstrated a capacity retention of approximately 90%, whereas the LiAlCl₄-based cell exhibited a faster degradation rate. The superior performance of the Li–In|Li₆PS₅Cl|2(Li_1.6AlCl_3.4S_0.6):TiS₂ cell is attributed to the synergistic benefits of enhanced ionic conductivity, electrochemical stability, and improved utilization of cathode active materials.⁴⁷

Conclusion

Developing inexpensive SEs using earth-abundant elements is imperative to reduce the cost of ASSBs for widespread applications in electric vehicles and consumer electronics. In this work, we synthesized a new lithium chalcohalide solid electrolyte, Li_1.6AlCl_3.4S_0.6, with a room-temperature ionic conductivity of 0.18 mS cm⁻¹. Structural characterization reveals that the enhanced ionic conductivity of Li_1.6AlCl_3.4S_0.6 strongly correlates with the formation of face- and edge-sharing octahedrally coordinated lithium sites, which often create low-energy conduction pathways. AIMD simulations using the refined structures reveal incorporating the optimal amount of S into LiAlCl₄ transitions it from a 1D to a 3D conductor. ^6,7Li MAS NMR combined with tracer exchange and relaxometry reveals increased ion mobility and participation of all Li⁺ sites in ion conduction. Li_1.6AlCl_3.4S_0.6 demonstrates good long-term cycling stability and rate performance in ASSBs, achieving a specific capacity of 180 mA h g⁻¹ at a fast charging rate of 2C in a Li–In|Li₆PS₅Cl|2(Li_1.6AlCl_3.4S_0.6):TiS₂ battery cell, compared to 37 mA h g⁻¹ in the LiAlCl₄-containing cell. The cost-effectiveness, combined with the demonstrated high performance, makes Li_1.6AlCl_3.4S_0.6 an excellent candidate as electrolytes for ASSBs.

Data availability

The data that supports this manuscript will be available upon request.

Author contributions

Tej Prasad Poudel: data curation: lead; formal analysis: lead; investigation: lead; methodology: lead; software: lead; validation: lead; visualization: lead; writing – original draft: lead. Ifeoluwa P. Oyekunle: data curation: equal; formal analysis: equal; investigation: equal; methodology: equal; software: equal; validation: equal; visualization: equal; writing – original draft: equal. Michael J. Deck: formal analysis: supporting; investigation: supporting; methodology: supporting; validation: supporting; visualization: supporting; writing – original draft: supporting. Yudan Chen: data curation: supporting; formal analysis: supporting; investigation: supporting; methodology: supporting; software: supporting; writing – original draft: supporting. Dewen Hou: data curation: supporting. Pawan K. Ojha: data curation: supporting; formal analysis: supporting; visualization: supporting; writing – review & editing: supporting. Bright O. Ogbolu: validation: supporting; writing – review & editing: supporting. Chen Huang: software: supporting. Hui Xiong: resources: supporting; supervision: supporting; validation: supporting. Yan-Yan Hu, PhD: conceptualization: lead; formal analysis: lead; funding acquisition: lead; methodology: lead; project administration: lead; resources: lead; supervision: lead; validation: lead; visualization: equal; writing – original draft: equal; writing – review & editing: lead.

Conflicts of interest

The authors declare no conflict of interest.

Acknowledgements

The authors acknowledge the support from the National Science Foundation under grant no. DMR-1847038 and DMR-2319151. Hou and Xiong thank the support from the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences program under Award Number DE-SC0019121. All solid-state NMR experiments were performed at the National High Magnetic Field Laboratory, which is supported by National Science Foundation Cooperative Agreement No. DMR-1644779 and DMR-2128556*. This research used resources of the Advanced Photon Source, a U.S. Department of Energy (DOE) Office of Science user facility operated for the DOE Office of Science by Argonne National Laboratory under Contract No. DE-AC02-06CH11357.

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Footnotes

† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d4sc07151d

‡ Authors contributed equally.

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